1. STANDARD ELECTRODE
POTENTIAL S
EXPLAIN THE MEASUREMENTS OF
THE STANDARD ELECTRODE
POTENTIALS

2. DEFINITION
• the 'potential' of a redox system to lose or
gain electrons when compared to the
standard hydrogen electrode - assigned a
value of 0 volts.
• Def of standard electrode potential- potential
diff between a standard hydrogen electrode
and a metal which is immersed in a solution
containing metal ions at 1 mol dm-3 conc. at
298 K and 1 atm pressure

3. • In any reduction-oxidation half equation the
electrons are gained by the species on the left
hand side:
Cu2+ + 2e Cu
• This is an equilibrium and so if a more powerful
reducing agent is allowed enters into electrical
contact with the above system it can force the
copper ions to accept electrons and push the
equilibrium to the right hand side.
• Conversely, if a weaker reducing agent is brought
into contact with the above equilibrium then the
copper can force it to accept electrons allowing
its own equilibrium to move to the left hand side.

4. • The electrode potential measures the tendency
of electrons to flow away from or towards a
redox equilibrium. They are always measured
with respect to the standard hydrogen electrode
(which is assigned a value of zero volts).
• Equilibrium redox systems with the reduced side
(usually a metal) more reactive than hydrogen
have a negative electrode potential, i.e. they can
lose electrons more easily than hydrogen.
• Equilibrium redox systems with the reduced side
less reactive than hydrogen have a positive
electrode potential, i.e. they can lose electrons
less easily than hydrogen.

5. EXAMPLE
• Zinc has a standard electrode potential of - 0.76 volts
• Consequently the equilibrium...
Zn Zn2+ + 2e
• has more of a tendency to move to the right hand side
than the equilibrium...
H2 2H+ + 2e
• Hence if the two equilibria are brought into electrical
contact using an external wire and a salt bridge, the
electrons will be pushed from the zinc equilibrium to
the hydrogen equilbrium with a force of - 0.76V (the
negative sign simply indicates the direction of flow -
from zinc to hydrogen ions)

6. • The two equations then may be summed
together to give the reaction occuring in the
whole cell.
• Zn Zn2+ + 2e
• 2H+ + 2e H2
• overall cell reaction
• Zn + 2H+ Zn2+ + H2

8. Representing the cell
• The whole cell can be represented by showing
the half cells in order of phase (solid, |
solution, |salt bridge | solution | solid)